General Chemistry Laboratory
The Decomposition of Hydrogen Peroxide
• To learn about Avogadro’s Hypothesis and the Ideal Gas Law.
• To learn about gas phase Stoichiometry.
• To learn about the chemistry of Peroxides.
In this laboratory exercise we will determine the percentage Hydrogen Peroxide present in a
commercially available solution by measuring the volume of Oxygen gas liberated when the
peroxide decomposes. Because the decomposition of dilute solutions of Hydrogen Peroxide is
relatively slow, a Ferric Chloride catalyst will be used to increase the reaction rate so that the
reaction goes to completion during the laboratory period.
In its pure form, Hydrogen Peroxide (H2O2) is a faintly bluish, syrupy liquid which boils at
150.2oC. It was first synthesized by the French chemist Louis Jaques Thenard in 1818 by
acidification of Barium Peroxide (BaO2) with Nitric Acid (HNO3). This process was supplanted
by an improved version in which the Barium Peroxide is initially treated with Hydrochloric Acid
(HCl), followed by addition of Sulfuric Acid (H2SO4) to precipitate the Barium Ion (Ba2+) as
Barium Sulfate (BaSO4). This leaves a relatively pure aqueous solution of Hydrogen Peroxide:
BaO2(s) + 2 HCl(aq) H2O2(aq) + BaCl2(aq) (Eq. 1)
BaCl2(aq) + H2SO4(aq) BaSO4(s) + 2 HCl(aq) (Eq. 2)
When exposed to sunlight or metallic impurities, Hydrogen Peroxide rapidly decomposes to
2 H2O2(aq) 2 H2O(l) + O2(g) (Eq, 3)
This is a key reaction of Hydrogen Peroxide.
Hydrogen Peroxide solutions (3-30%) are used for bleaching (pulp, paper, straw, leather, hair,
etc.) and to treat wounds. Its value as an antiseptic is low, but the evolution of Oxygen when it
comes into contact with clotted blood helps to loosen dirt and assists in cleansing a wound.
At higher concentrations (70-98%), the decomposition of the peroxide is accompanied by the
evolution of enough heat to convert the Water to Steam. In this fashion, Hydrogen Peroxide is
used as a monopropellant in rocket engines; the peroxide is passed over a Silver mesh which
catalyzes the decomposition and the resulting gaseous H2O and O2 products are ejected through a
nozzle at high velocity propelling the rocket forward. Concentrated Hydrogen Peroxide can also
be used as an oxidant with organic compounds, such as kerosene, in a bipropellant rocket engine.
The German V2 Rocket of WWII used this design.
We will leverage the decomposition reaction (Eq. 3) to determine the concentration of Hydrogen
Peroxide in products typically sold in the supermarket. The reaction stoichiometry allows us to
use the measured amount of Oxygen (#moles O2) produced to determine the amount of peroxide
(#moles H2O2) initially present in the solution:
# moles H2O2 = x # moles O2 (Eq. 4)
The number of moles Oxygen produced (n) by this reaction can be determined by simply
measuring the volume of Oxygen gas (V) generated; the volume of a gas being related to the
number of moles of its constituents by Avogadro’s Hypothesis.
In 1811, the Italian physicist Amadeo Avogadro advanced the hypothesis that equal volumes of
all gases (V), measured under the same conditions of temperature (T) and pressure (P), contain
the same number of molecules. As we count molecules by the mole (n), we have;
V ~ n (Eq. 5)
At the time of its statement, this hypothesis made little or no impact on Avogadro's
contemporaries, in large part because his paper in the Journal de Physique contained little
experimental data. Avogadro's idea lay fallow for nearly half a century until after his death his
student Stanislo Cannizzaro revived the hypothesis and showed that it resolved a number of
conflicts concerning the atomic weights of the elements. Since that time, considerable evidence
has accumulated to demonstrate the validity of Avogadro’s Hypothesis. Taken together with the
other historical gas laws (Boyle’s, Charles’, Gay-Lusaacs’), Avogadro’s Hypothesis forms the
basis of the Ideal Gas Law (PV = nRT). Thus, the number of moles of a gas can be determined
by measuring its volume at a given temperature and pressure:
As a word of caution, it must be remembered the Ideal Gas Law is a Limiting Law; it is strictly
true only in the limit of zero pressure (Why?):
At atmospheric pressures (P ~ 760 mmHg), application of the Ideal Gas Law can give results
which are in error by as much as 10% or more.
As an Example of the application of this procedure, consider a different decomposition reaction in
which a single reaction product is gaseous:
2 KClO3(s) 2 KCl(s) + 3 O2(g)
Suppose we collect 638mL of Oxygen gas at 128oC and 752mmHg. What mass of Potassium
Chlorate (KClO3) did we start with?
First, we convert the units of each measured quantity to those consistent with the units of the
Universal Gas Constant R = 0.08206 L atm / K mole:
V = 0.638 L
T = 401K
P = 0.989 atm
Second, we use the Ideal Gas Law (containing Avogadro’s Hypothesis) to determine the number
of moles Oxygen produced:
n = (0.989 atm) (0.638 L) / (0.08206 L atm / K mole) (401K)
= 0.0192 mole O2
Finally, the reaction stoichiometry is applied to determine the amount Potassium Chlorate used:
# moles KClO3 = (2 moles KClO3 / 3 moles O2) x (0.0192 moles O2)
= 0.0128 mole KClO3
#g = (0.0128 mole KClO3) x (122.55 g/mole)
Of course, in this case we could have simply weighed the amount of KClO3 we started with
instead of using this round-about method. However, in our case we cannot do this because the
Hydrogen Peroxide is dissolved in an aqueous solution.
We must arrange our experiment such that the Oxygen gas produced is trapped so that its volume
can be determined. We will do this by attaching the reaction vessel to a tank filled with Water
which is itself arranged such that the Water will be pushed out as Oxygen fills the tank.
Using this arrangement, we can determine the volume Oxygen gas produced by simply
measuring the volume of the displaced Water. This experimental arrangement presents us with
one minor complication, however; the gas in the Collection Tank is composed of generated
Oxygen and Water Vapor in equilibrium with the liquid Water. Fortunately, the presence of the
Water Vapor can be accounted for using Dalton’s Law of Partial Pressures. According to
Pgas = Σ Pi (Eq. 8)
where Pi represents the Partial Pressure of each gas in the mixture. For the present case, this
Pgas = PO2 + PH2O (Eq. 9)
where PH2O is the Vapor Pressure of the Water. Tabulated values of the Vapor Pressure of Water
as a function of temperature are provided in the Appendix. This means the Partial Pressure of
the Oxygen gas, PO2, can be determined by measuring the Pressure of the gas in the Collection
Tank, Pgas, and subtracting the Vapor Pressure of Water.
We have one last hurdle to overcome; how do we measure Pgas? We employ a trick! If the
surface of the Displaced Water is arranged so that it is leveled with the surface of the Water in
the Collection Tank, then the pressure of the gas in the Collection Tank is equivalent to the
atmospheric pressure, Patm:
Pgas = Patm (Eq. 10)
This is important because the atmospheric pressure can be easily measured using a relatively
simple barometer. Thus, the partial pressure of Oxygen in the Collection Tank can be
PO2 = Patm - PH2O (Eq. 11)
Finally, we should consider the role of the Ferric Chloride (FeCl3) catalyst in the decomposition
reaction. In the absence of a catalyst, Hydrogen Peroxide is quite stable, decomposing only very
slowly. (Otherwise you would not be able to buy it in the supermarket.) However, when any of
a number of catalysts is present, the decomposition reaction proceeds very rapidly. And, many
of these catalysts take advantage of the fact that this reaction is a Disproportionation; the
Hydrogen Peroxide is both Oxidized and Reduced. Peroxides are unique in that the Oxygen
exists in a -1 oxidation state (O-1), which lies between the usual states of O0 and O-2. Thus, the
Hydrogen Peroxide can disproportionate to both O0 and O-2 according to:
O-1 + e- O-2 Reduction
2 H2O2 2 H2O + O2
O-1 O0 + e- Oxidation
This, coupled with the fact that Iron can exist in two different oxidation states, Fe2+ (Ferrous) and
Fe3+ (Ferric), allows the catalyst to break the reaction into two different redox steps, each of
which has a lower energy barrier to completion than the uncatalyzed reaction:
H2O2(aq) + 2 Fe3+(aq) O2(g) + 2 Fe2+(aq) + 2 H+(aq) (Eq. 12)
H2O2(aq) + 2 Fe2+(aq) + 2 H+(aq) 2 H2(l) + 2 Fe3+(aq) (Eq. 13)
Note the first step in the catalyzed reaction involves reduction of the Ferric Ion (Fe3+) to the
Ferrous Ion (Fe2+), which is then re-oxidized to Ferric Ion in the second step. Hence, on net, the
catalyst is not consumed during the course of the decomposition.
Transition Metals like Iron (Fe) are not the only possible catalysts for this reaction. Because
Hydrogen Peroxide is harmful biologically, a number of enzymes, large protein molecules that
catalyze biologically important reactions, also catalyze this reaction. Catalase, present in the
blood and liver of mammals, is an example of such an enzyme. The active site of Catalase, the
point at which the reaction takes place, contains an Fe3+ ion attached to a Heme group embedded
in the protein. Catalase most likely catalyzes the decomposition of Hydrogen Peroxide via a
mechanism that is very similar to that used by Ferric Chloride. (Take Home Expt Idea: The next
time you have liver for dinner, try grinding up a bit of the raw liver and add it to a solution of
Hydrogen Peroxide. You should see an immediate fizzing due to the production of Oxygen gas.
The liver must be ground in order to break apart the cells and release the Catalase.)
In summary, we will measure the volume of Oxygen gas produced, at a given temperature and
pressure, as a result of the Ferric Chloride catalyzed decomposition of Hydrogen Peroxide in a
commercial solution. The measured volume will allow us to determine the number of moles
Oxygen produced. The reaction stoichiometry will then be used to determine the number of
moles Hydrogen Peroxide in the starting solution. This information can then be used to
determine the percentage Hydrogen Peroxide in the solution.
1. A sample of Oxygen gas is collected over Water at 28oC in an experimental set-up similar
to that employed in this laboratory exercise. The atmospheric pressure was measured using
a barometer and found to be 752.7 mmHg. What is the partial pressure of the Oxygen gas
in the Collection Tank?
2. How can we improve the measurement of the volume of Water displaced? (Hint: Which can
be measured more precisely, the volume or mass of a substance?)
3. When metallic Zinc is placed in an acidic solution, Hydrogen gas is produced:
Zn(s) + H2SO4(aq) ZnSO4(aq) + H2(g)
159 mL of Hydrogen gas (dry) is collected at 24oC and 738.4 mmHg.
a) How many moles of Hydrogen gas was produced by this reaction?
b) How many grams of Zinc (Zn) were used initially?
4. A good barometer can easily measure the atmospheric pressure with an error of less than
0.1%. Suppose a manufacturer has developed a better barometer that can reduce this error
to less than 0.01%, but this barometer is very expensive. Would it be worth the expense to
employ this new barometer in this laboratory exercise? Explain briefly.
1. Assemble the Reaction Apparatus as diagramed above. Use a 500 mL Erlenmeyer flask for
the Reaction Vessel. Use an 800mL Beaker to collect the Displaced Water.
2. Clean and dry the Beaker used to collect the Displaced Water.
3. Measure out 30.0mL of commercial Hydrogen Peroxide solution in a graduated cylinder
and pour it into the Reaction Vessel. (Record the exact volume used, the Brand and the
reported percentage Hydrogen Peroxide.)
4. Measure ~6mL of the catalyst (FeCl3) into a small test tube. Make sure none of it
contaminates the outside of the test tube. Slide the test tube into the flask so that it does not
spill into the H2O2 solution.
5. Your instructor will demonstrate the following step! Fill the Oxygen Collection Tank
until its almost full of Water. With the pinch clamp on the delivery tube open, and the
Reaction Vessel unstoppered, establish a siphon from the Collection Tank to the Displaced
Water beaker. (A simple Pipette Bulb can be used to “push” the Water through the system
from the Reaction Vessel side of the Collection Tank.) Be sure and remove all air bubbles
in the delivery tube. Collect ~200 mL of Water in the Beaker, enough to completely cover
the delivery tube. Now, close the pinch clamp.
6. Connect the Oxygen delivery tube to the Reaction Vessel and make sure the Vessel is
7. Open the pinch clamp and level the Water in the Collection Tank with the Water in the
Displaced Water beaker. (Recall, this is done to ensure the gas’s pressure Pgas in the
closed system is the same as the atmospheric pressure Patm.) Close the pinch clamp when
this is the case.
8. Carefully lower the Beaker and remove the delivery tube from the beaker so as to not lose
the siphon. (If you do lose the siphon, you will have to start the leveling procedure all over
again.) Pour out all the Water in the Displaced Water beaker, dry the beaker and replace it.
Again, this must be done such that your siphon is not lost. Open the pinch clamp. If done
correctly, a little Water will dribble into the Displaced Water beaker as the siphon tries to
re-establish itself. However, the Water flow should cease almost immediately. If it does
not, you have a leak in the system. If this is the case, check to make sure all the
connections are tight and re-level the system. (Why do we not need to subtract this bit of
Water from the volume of Water collected after the evolution of the Oxygen gas is
9. Tip the Reaction Vessel so that the catalyst flows into the Hydrogen Peroxide solution.
Swirl the mixture. You should observe the production of a gas after a few moments.
About half-way through the reaction (~7 minutes) swirl the system again to make sure the
reaction goes to completion.
10. When the reaction is complete (~15minutes), the level of the Water in the Collection Tank
will remain unchanged. At this point, again level the water in the two containers. Close
the pinch clamp.
11. Record the temperature of the Water in the Displaced Water beaker. This will be assumed
to be approximately the temperature of the gas in the Collection Jar.
12. Measure the volume of Water displaced using a large graduated cylinder.
13. Record the barometric pressure. Your laboratory instructor will show you how to correctly
read the barometer.
1. Determine the Partial Pressure of the Oxygen gas collected, PO2.
2. Determine the number of moles Oxygen gas collected.
3. Determine the number of moles Hydrogen Peroxide present in the initial solution. Convert
this to the mass of Hydrogen Peroxide.
4. Calculate the mass of the initial Solution. The density of a dilute Hydrogen Peroxide
solution near Room Temperature is 1.03 g/mL.
5. Determine the percentage Hydrogen Peroxide in the commercial solution.
6. What is the percentage difference between your determination and that reported by the
7. Comment on which measurement or procedure has the largest error. How might you
improve this measurement?
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Post Lab Questions
1. Hydrazine (NH2NH2) is considered the Group 5A cousin of Hydrogen Peroxide and is also
used as a rocket fuel, particularly as a fuel in ICBM’s. Draw the Lewis Structures of
Hydrazine and Hydrogen Peroxide.
2. The decomposition of KClO3 is another example of a disproportionation:
2 KClO3(s) 2 KCl(s) + 3 O2(g)
Determine the oxidation states for the Chlorine and Oxygen in each of the above species
and identify the Oxidation and Reduction half-reactions.
3. We have indicated the oxidation state of the Oxygen in H2O2 is O-1, midway between the
extremes of O0 and O-2. Hence, Hydrogen Peroxide can act as an Oxidizing Agent and a
Reducing Agent. In practice, Hydrogen Peroxide is a powerful Oxidizing Agent and only a
weak Reducing Agent. In the following reaction it is acting as an Oxidizing Agent:
H2O2(aq) + 2 I-(aq) + 2 H+(aq) 2 H2O + I2(aq)
Determine the oxidation States for the Oxygen and Iodine in each of the above species and
identify the Oxidation and Reduction half-reactions.
4. Which ion acts as a “Spectator” in (Eq. 1) and (Eq. 2)
5. When collecting Oxygen gas over Water, we have to correct the gas pressure by subtracting
out the vapor pressure of Water because we know some Water vaporizes into the gas phase.
We also know Oxygen gas will dissolve, to some extent, in liquid Water. Why do we not
have to correct our results for the solubility of Oxygen in the Water in the Collection Tank?
P a g e | 11
Appendix - Some Physical Properties of Water
Temp [oC] Vapor Press [mmHg] Density [g/mL]
0 4.6 0.99984
5 6.5 0.99997
10 9.2 0.99970
11 9.8 0.99960
12 10.5 0.99950
13 11.2 0.99938
14 12.0 0.99925
15 12.8 0.99910
16 13.6 0.99895
18 15.5 0.99860
20 17.5 0.99821
22 19.8 0.99777
24 22.4 0.99730
26 25.2 0.99679
28 28.3 0.99624
30 31.8 0.99565
35 42.2 0.99403
40 55.3 0.99222
45 71.9 0.99022
50 92.5 0.98803
55 118.0 0.98570
60 149.4 0.98320
65 187.5 0.98056
70 233.7 0.97778
75 289.1 0.97485
80 355.1 0.97182
85 433.6 0.96862
90 525.8 0.96535
95 633.9 0.96190
100 760.0 0.95840
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A Reaction of Sodium Peroxide
In this short exercise, we will treat a mixture of Sodium Peroxide (Na2O2) and Sulfur (S) with
Water. This results in the production of aqueous Hydrogen Peroxide which then oxidizes the
Na2O2(s) + 2 H2O(l) 2 H2O2(aq) + 2 NaOH(aq) (Eq. A1)
2 H2O2(aq) + S(s) 2 H2O(l) + SO2(g) (Eq. A2)
We can see the Hydrogen Peroxide is acting as an Oxidizing Agent in reaction (Eq. A2) by
explicitly assigning appropriate oxidation numbers to the relevant species:
O-1 + e- O-2 Reduciton
2 H2O2 + S 2 H2O + SO2
S0 S4+ + 4 e- Oxidation
The steps in this procedure must be followed very carefully. The Sodium Peroxide can
detonate prematurely if not handled with caution. If you are unsure of what to do, have
your instructor help you do this experiment. Do this experiment in a Fume Hood!
1. Wearing gloves, obtain ~1.0g Sodium Peroxide (Na2O2) that has been previously weighed
out. Sodium Peroxide is a strong Oxidizing Agent. Avoid getting on your skin or
clothing. It is also Water sensitive, so do not place your sample into anything that is
2. Also obtain ~0.1g of Sulfur (S) that has also been previously weighed out. In a very dry
evaporating dish gently mix these solids together with your spatula to form a uniform
3. Form the powder into a small pile in the middle of the evaporating dish. Make a slight
depression in the top of the pile.
4. Add 1-2 drops of water to the depression and quickly step back. Observe the results.